Inorganic Chemistry-1 BCH1T01

Syllabus topics (official RTMNU)

(A) Atomic structure

• Bohr model, Idea of de-Broglie matter Waves, Heisenberg's uncertainty principle.
• Schrodinger wave equation, significance of ψ and ψ2, Quantum numbers, Concept of atomic orbital, Radial and angular wave functions and probability distribution curves for 1s, 2s, 2p, 3s, 3p and 3d orbitals, shapes of s, p and d orbitals, Aufbau principle, Bohr-Bury rule, Pauli's exclusion principle and Hund's rule of maximum multiplicity.
• Principle of extra stability.
• Electronic Configuration of elements and ions.

(B) Periodic Properties

• Atomic and ionic radii, ionization energy, electron affinity and electronegativity- Definition, trends in periodic table.
• Factors affecting ionization potential.
• Pauling's, Mulliken's and Allred-Rochow scale of electronegativity.
• Effective nuclear charge and Slater's rule with some numericals.

(A) Atomic structure: Bohr model to quantum model

Bohr model (1913)

• Postulate 1: electrons revolve around the nucleus in fixed circular orbits called stationary states or energy levels (n = 1, 2, 3…). In these orbits they do not radiate energy.
• Postulate 2: angular momentum is quantised: mvr = nh/2π.
• Postulate 3: energy is absorbed or emitted only when an electron jumps from one orbit to another: ΔE = E₂ − E₁ = hν.
• Results for hydrogen-like species: rₙ = 0.529 n²/Z Å and Eₙ = −13.6 Z²/n² eV (−2.18 × 10⁻¹⁸ Z²/n² J).
• Limitations: works only for one-electron species (H, He⁺, Li²⁺); cannot explain spectra of multi-electron atoms, fine structure, Zeeman and Stark effects; goes against the uncertainty principle (fixed path).

de Broglie matter waves (1924)

Every moving particle has a wave nature. Wavelength λ = h/mv = h/p. For an electron in a Bohr orbit, a stable orbit holds a whole number of waves: 2πr = nλ, which gives back mvr = nh/2π. The wave nature is noticeable only for very small particles (electrons), not for a cricket ball.

Heisenberg's uncertainty principle (1927)

It is impossible to measure the exact position and exact momentum of a microscopic particle at the same time: Δx · Δp ≥ h/4π. So we cannot speak of a fixed orbit; we can only speak of the probability of finding the electron.

Schrödinger wave equation and orbitals

Schrödinger treated the electron as a 3-D standing wave:

∂²ψ/∂x² + ∂²ψ/∂y² + ∂²ψ/∂z² + (8π²m/h²)(E − V)ψ = 0 (short form Ĥψ = Eψ)

• ψ (wave function): amplitude of the electron wave; it can be +, − or zero and has no physical meaning by itself.
• ψ²: probability density, i.e. probability of finding the electron per unit volume at a point (Born interpretation). It is always positive.
• Acceptable ψ must be single-valued, finite, continuous, and normalised (∫ψ² dτ = 1).
• Atomic orbital: the region around the nucleus where the probability of finding the electron is maximum (about 90–95 %). It is a solution ψ of the wave equation with definite n, l, m.

Quantum numbers

Radial and angular wave functions; probability curves

ψ is written as ψ(r, θ, φ) = R(r) · Y(θ, φ). R(r) is the radial part (depends on n and l) and Y is the angular part (depends on l and m; decides shape).

• Radial probability distribution = 4πr²R² plotted against r. It shows the probability of finding the electron in a thin spherical shell at distance r.
• Radial nodes = n − l − 1; angular nodes = l; total nodes = n − 1.
• 1s: one peak, maximum at 0.529 Å (Bohr radius), no node.
• 2s: two peaks (small inner, large outer), one radial node. 3s: three peaks, two nodes.
• 2p: one peak, no radial node. 3p: two peaks, one radial node.
• 3d: one peak, no radial node (n − l − 1 = 0), lies closer to nucleus than 3s/3p main peaks.
• The small inner peaks of s orbitals show penetration towards the nucleus, so for the same n the order of energy is s < p < d.

Shapes of orbitals

• s: spherical, non-directional.
• p (pₓ, p_y, p_z): dumb-bell shaped, two lobes with a nodal plane through the nucleus.
• d: five orbitals. d_xy, d_yz, d_zx have four lobes between the axes; d_x²−y² has four lobes along the x and y axes; d_z² has two lobes along z with a ring (doughnut) in the xy plane.

Rules for filling electrons

• Aufbau principle: orbitals are filled in order of increasing energy. (n + l) rule: lower n + l fills first; if equal, lower n fills first. Order: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p…
• Bohr–Bury rule: maximum electrons in a shell = 2n²; the outermost shell cannot have more than 8 and the penultimate not more than 18 electrons.
• Pauli's exclusion principle: no two electrons in an atom can have all four quantum numbers the same. So one orbital holds maximum 2 electrons with opposite spins.
• Hund's rule of maximum multiplicity: in degenerate orbitals, pairing starts only after each orbital has one electron with parallel spin. Example N: 2p³ = ↑ ↑ ↑ (three unpaired electrons).

Extra stability of half-filled and completely filled subshells

Due to symmetrical distribution and maximum exchange energy, d⁵, d¹⁰, f⁷, f¹⁴ are extra stable. Examples: Cr (24) = [Ar] 3d⁵ 4s¹ (not 3d⁴ 4s²); Cu (29) = [Ar] 3d¹⁰ 4s¹.

Electronic configuration of ions

For cations of transition metals, electrons are removed first from the outer 4s, then 3d. Fe (26) = [Ar] 3d⁶ 4s²; Fe²⁺ = [Ar] 3d⁶; Fe³⁺ = [Ar] 3d⁵ (extra stable, so Fe³⁺ is more stable than Fe²⁺). Anions add electrons: Cl⁻ = [Ne] 3s² 3p⁶.

(B) Periodic properties

Factors affecting ionisation potential

• Atomic size: larger size → lower IE.
• Nuclear charge: higher effective nuclear charge → higher IE.
• Screening (shielding) by inner electrons → lower IE.
• Penetration: s > p > d > f; an s electron is harder to remove.
• Stable configuration: half-filled/fully filled shells have high IE. So IE of N > O and Be > B; noble gases highest.

Electronegativity scales

• Pauling scale: from bond energies. Δ = E(A–B) − √[E(A–A) × E(B–B)]; χA − χB = 0.102 √Δ (Δ in kJ mol⁻¹) [or 0.208 √Δ with Δ in kcal mol⁻¹]. F = 4.0 (highest).
• Mulliken scale: average of IE and EA: χ = (IE + EA)/2. With values in eV, χ(Pauling) ≈ χ(Mulliken)/2.8.
• Allred–Rochow scale: EN is the electrostatic force exerted by the nucleus on a valence electron at the covalent radius: χ = 0.359 Z_eff/r² + 0.744 (r = covalent radius in Å).

Effective nuclear charge and Slater's rules

Z_eff = Z − σ, where σ = screening constant. Slater's rules: write groups (1s)(2s,2p)(3s,3p)(3d)(4s,4p)(4d)…

• Electrons to the right of the chosen electron: 0.
• For an ns/np electron: others in the same group 0.35 each (0.30 for 1s); (n−1) shell 0.85 each; (n−2) and lower 1.00 each.
• For a d or f electron: same group 0.35; all electrons to the left 1.00.

Likely exam questions

Long answer

1. Explain the significance of ψ and ψ². Draw and discuss the radial probability distribution curves for 1s, 2s, 2p, 3s, 3p and 3d orbitals.
2. Define electronegativity. Describe the Pauling, Mulliken and Allred–Rochow scales of electronegativity.

Short answer

1. State Heisenberg's uncertainty principle and give its mathematical form.
2. State Aufbau principle, Pauli's exclusion principle and Hund's rule with one example each.
3. Calculate the effective nuclear charge for the 4s electron of potassium using Slater's rules.

Syllabus topics (official RTMNU)

(A) Covalent Bond

• Valency Bond Theory, Formation of Hydrogen molecule with potential energy diagram with all improvements, Limitations of VBT, Bond energy, bond length, Bond order, Bond angle.
• Various types of hybridization and shape of inorganic molecules [BeF2, BCl3, CH4, NH3, H2O, PCl5, SF6 and IF7].
• VSEPR Theory: Rules/postulates and their applications to various common molecules and ions (NH3, ClF3, H2O, SF4, H3O+, NH4+, ICl2- etc)

(B) Ionic solids

• Close packing in ionic solids [Square, Hexagonal, Cubic, BCC and FCC], Radius ratio rule and its limitations, Ionic structures with respect to NaCl and CsCl, Lattice energy and Born- Haber cycle.
• Solvation energy and solubility of ionic solids, polarizing power and polarizability of ions, Fajan’s rules.

(A) Covalent bond: Valence Bond Theory (VBT)

A covalent bond is formed by sharing of a pair of electrons between two atoms. VBT (Heitler and London, 1927; extended by Pauling and Slater) says a bond forms when half-filled atomic orbitals of two atoms overlap and the two electrons pair up with opposite spins. Greater overlap = stronger bond.

Formation of H₂ molecule – potential energy diagram

• Two H atoms far apart: no interaction, energy taken as zero.
• As they come closer, attraction (nucleus of one – electron of other) is more than repulsion, so energy falls.
• At r₀ = 0.74 Å (74 pm) energy is minimum: this is the bond length. Energy released = bond energy ≈ 436 kJ mol⁻¹ (458 kJ mol⁻¹ calculated after all corrections).
• Closer than r₀, nucleus–nucleus repulsion rises sharply and energy increases.

Diagram in words: y-axis = potential energy, x-axis = internuclear distance. The curve comes down from zero, makes a deep minimum at 0.74 Å, then rises steeply at short distance. The curve for parallel spins (repulsive state) always stays above zero with no minimum.

Limitations of VBT

• Cannot explain paramagnetism of O₂.
• Cannot explain bonding in electron-deficient molecules (B₂H₆) and odd-electron species (H₂⁺).
• Does not explain colour and spectra of complexes; gives only qualitative results.

Bond parameters

• Bond length: equilibrium distance between nuclei of two bonded atoms (C–C 154 pm, C=C 134 pm, C≡C 120 pm).
• Bond energy: energy required to break one mole of bonds in gaseous molecules.
• Bond order: number of bonds between two atoms. Higher bond order → shorter and stronger bond.
• Bond angle: angle between two bonds around a central atom (H₂O 104.5°).

Hybridisation and shapes

Hybridisation is mixing of atomic orbitals of nearly equal energy on the same atom to give the same number of equivalent hybrid orbitals.

VSEPR theory

Postulates (Sidgwick–Powell, Gillespie–Nyholm):

1. Shape depends on the total number of electron pairs (bond pairs + lone pairs) around the central atom.
2. Electron pairs repel each other and arrange themselves as far apart as possible.
3. Order of repulsion: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair.
4. Multiple bonds are treated as a single bonding region; they repel more than single bonds.
5. Repulsion decreases as electronegativity of the surrounding atom increases.

(B) Ionic solids

Close packing

• Square packing (2-D): each sphere touches 4 others; space use 52.4 % (simple cubic in 3-D, coordination number 6).
• Hexagonal packing (2-D): each sphere touches 6; layers stacked ABAB… give hcp, ABCABC… give ccp (= fcc). Both have coordination number 12 and 74 % packing.
• BCC: coordination number 8, 68 % packing, 2 atoms per unit cell.
• FCC: coordination number 12, 74 % packing, 4 atoms per unit cell.

Radius ratio rule

The ratio r⁺/r⁻ decides coordination number and structure.

Limitations: ions are not hard spheres; radii are not exact; covalent character (polarisation) is ignored; some compounds do not obey it (e.g. LiI).

NaCl and CsCl structures

• NaCl: fcc lattice of Cl⁻ with Na⁺ in all octahedral holes; 6:6 coordination; 4 NaCl units per unit cell; r⁺/r⁻ ≈ 0.52.
• CsCl: simple cubic Cl⁻ with Cs⁺ at the body centre; 8:8 coordination; 1 CsCl unit per cell; r⁺/r⁻ ≈ 0.93.

Lattice energy and Born–Haber cycle

Lattice energy (U) is the energy released when one mole of an ionic solid is formed from its gaseous ions: Na⁺(g) + Cl⁻(g) → NaCl(s). Born–Haber cycle (Hess's law) for NaCl:

ΔHf = S + ½D + IE + EA + U

where S = sublimation of Na (+108), ½D = half dissociation of Cl₂ (+121), IE of Na (+496), EA of Cl (−349), ΔHf = −411 kJ mol⁻¹. So U = −411 − 108 − 121 − 496 + 349 = −787 kJ mol⁻¹.

Solvation energy and solubility

An ionic solid dissolves if hydration (solvation) energy ≥ lattice energy. High dielectric constant of water reduces attraction between ions. Example: LiF has very high lattice energy, so it is less soluble than LiCl.

Polarisation and Fajans' rules

Distortion of the anion's electron cloud by a cation is polarisation; it gives covalent character to an ionic bond. Covalent character increases when:

• cation is small and highly charged (high polarising power; e.g. Al³⁺ > Mg²⁺ > Na⁺);
• anion is large and highly charged (high polarisability; I⁻ > Br⁻ > Cl⁻ > F⁻);
• cation has pseudo-noble-gas (18-electron) configuration (Cu⁺, Ag⁺) rather than noble-gas configuration.

Examples: AlCl₃ is more covalent than NaCl; AgCl is less soluble in water than NaCl; LiCl is more covalent than KCl.

Likely exam questions

Long answer

1. Explain the formation of H₂ molecule on the basis of valence bond theory with a potential energy diagram. Give the limitations of VBT.
2. State the postulates of VSEPR theory. Explain the shapes of ClF₃, SF₄, H₃O⁺ and ICl₂⁻.

Short answer

1. State the radius ratio rule and give its limitations.
2. Draw the Born–Haber cycle for NaCl and write the expression for lattice energy.
3. State Fajans' rules with suitable examples.

Syllabus topics (official RTMNU)

(A) s-block elements

• Electronic configuration, atomic and ionic radii, Ionisation potential, Reducing properties and Metallic Properties.
• Diagonal Relationships (Li-Mg).
• Metallic Bond – Free electron, Valence Bond and Band theory of metallic bond.
• Properties of conductors, insulators and semiconductors based on Band theory.
• Hydrogen bonding - Classification and effect of Hydrogen bonding on viscosity, solubility, melting point and boiling point.

(B) p-block elements

• Introduction to p-block elements, comparative study of groups 15, 16 and 17 with respect to atomic and ionic radii, ionisation potential, electron affinity, electronegativity, redox properties, oxidation state.
• Diagonal relationship (B-Si).

(A) s-Block elements

Group 1 (alkali metals: Li, Na, K, Rb, Cs, Fr) and Group 2 (alkaline earth metals: Be, Mg, Ca, Sr, Ba, Ra). The last electron enters the s orbital.

Diagonal relationship (Li–Mg)

Li resembles Mg (diagonal in the periodic table) because they have similar ionic size and polarising power (charge/radius). Points of similarity:

• Both are harder and have higher m.p. than other members of their groups.
• Both burn in air to give normal oxide (Li₂O, MgO) and also form nitrides (Li₃N, Mg₃N₂).
• Their carbonates decompose on heating: Li₂CO₃ → Li₂O + CO₂; MgCO₃ → MgO + CO₂.
• Their fluorides, carbonates and phosphates are sparingly soluble in water.
• Their chlorides are deliquescent, covalent and soluble in ethanol; they crystallise as hydrates (LiCl·2H₂O, MgCl₂·6H₂O).
• Nitrates on heating give NO₂ and O₂: 4LiNO₃ → 2Li₂O + 4NO₂ + O₂.

Metallic bond

1. Free electron (electron sea) theory – Drude and Lorentz

A metal is a lattice of positive kernels (ions) placed in a sea of mobile valence electrons. The attraction between kernels and the electron sea is the metallic bond. It explains electrical and thermal conductivity, lustre (electrons absorb and re-emit light), malleability and ductility (layers slide without breaking bonds). It cannot explain why some solids are semiconductors or insulators.

2. Valence bond (resonance) theory – Pauling

Metallic bonding is resonance among many covalent and ionic structures. Example Li: Li–Li bonds resonate between many positions (Li⁺ Li⁻ forms also). The vacant orbital (metallic orbital) allows electron movement. It explains bonding but not the conductivity values or lustre well.

3. Band theory (molecular orbital theory)

• When N atoms combine, N atomic orbitals give N molecular orbitals very close in energy – a continuous band.
• Valence band: band containing valence electrons. Conduction band: next higher band, empty or partly filled.
• Forbidden gap (Eg): energy gap between the two bands.

Doping: adding group 15 (P, As) to Si gives n-type (extra electrons); adding group 13 (B, Ga) gives p-type (positive holes).

Hydrogen bonding

A weak electrostatic attraction between an H atom covalently bonded to a highly electronegative atom (F, O, N) and another electronegative atom: X–H····Y. Strength ≈ 10–40 kJ mol⁻¹.

• Intermolecular H-bond: between different molecules, e.g. H₂O, HF (zig-zag chains), alcohols, p-nitrophenol. Causes association.
• Intramolecular H-bond (chelation): within the same molecule, e.g. o-nitrophenol, salicylaldehyde. Causes ring formation, no association.

Effects of H-bonding

• Melting and boiling point: intermolecular H-bonding raises them. H₂O (100 °C) vs H₂S (−60 °C); HF has higher b.p. than HCl. o-Nitrophenol (intramolecular) is more volatile and steam-volatile, p-nitrophenol is not.
• Solubility: compounds forming H-bonds with water are soluble (ethanol, NH₃, sugars, lower amines). o-Nitrophenol is less soluble than p-nitrophenol.
• Viscosity: more H-bonds → higher viscosity. Glycerol (3 –OH) > ethylene glycol (2 –OH) > ethanol.
• Other: ice is less dense than water (open cage structure); KHF₂ exists (HF₂⁻ ion) but KHCl₂ does not.

(B) p-Block elements: groups 15, 16 and 17

Diagonal relationship (B–Si)

• Both are non-metals/metalloids with high m.p.; poor conductors.
• Both form acidic oxides: B₂O₃ and SiO₂, which give borates and silicates with bases.
• Both form many volatile, flammable hydrides (boranes, silanes) that are hydrolysed by water.
• Halides (BCl₃, SiCl₄) are covalent and hydrolysed: SiCl₄ + 4H₂O → Si(OH)₄ + 4HCl.
• Both form carbides (B₄C, SiC) that are very hard.

Likely exam questions

Long answer

1. Discuss the band theory of metallic bonding. On its basis explain the properties of conductors, insulators and semiconductors.
2. What is hydrogen bonding? Classify it with examples and explain its effect on melting point, boiling point, solubility and viscosity.

Short answer

1. Explain the diagonal relationship between Li and Mg with any four points.
2. Compare groups 15, 16 and 17 with respect to electron affinity and oxidation states.
3. Explain the free electron theory of metallic bond.

Syllabus topics (official RTMNU)

Hydrides

• Comparative study with respect to structure of NH3, PH3, AsH3 and SbH3.
• Structure and bonding of diborane, structure of borazine.
• Allotropes of Carbon [Graphite, Diamond and Fullerene], Carbides: Classification and uses, Silicates: classification, Nitrides of sulphur: Structure of S4N4, Fluorocarbon and its uses.
• Oxides of Phosphorus: Structure and bonding in P2O3 and P2O5.
• Oxyacids of Phosphorous: Structures of H3PO2, H3PO3, H3PO4, H4P2O6 and H4P2O7
• Peroxyacids of Sulphur: Preparation and structure of Caro's and Marshall's acids.
• Interhalogens and Polyhalides: Preparation, properties and structure of Interhalogen compounds.
• Polyhalides- Classification, Structure of I3-, I5-, ICl4-.

Hydrides of group 15: NH₃, PH₃, AsH₃, SbH₃

• All are pyramidal with one lone pair on the central atom.
• Bond angle decreases down the group: the central atom becomes larger and less electronegative, bond pairs move away from it, bond-pair repulsion falls and lone pair pushes them closer. Heavier atoms use nearly pure p orbitals (angle near 90°).
• Basic strength: NH₃ > PH₃ > AsH₃ > SbH₃ (lone pair less available as size increases).
• Thermal stability decreases and reducing power increases down the group. B.p.: NH₃ high due to H-bonding; then PH₃ < AsH₃ < SbH₃.

Diborane (B₂H₆) and borazine

• Preparation: 4BF₃ + 3LiAlH₄ → 2B₂H₆ + 3LiF + 3AlF₃ (in ether).
• Diborane is electron-deficient: 12 valence electrons, but 8 bonds would need 16.
• Structure: each B is sp³ hybridised. Four terminal B–H bonds (normal 2-centre-2-electron bonds, 119 pm) lie in one plane. Two bridging H atoms lie above and below this plane, forming two 3-centre-2-electron (banana) bonds B–H–B (133 pm).
• Terminal H–B–H angle ≈ 122°; bridge H–B–H angle ≈ 97°.

Borazine, B₃N₃H₆ ("inorganic benzene")

• Preparation: 3B₂H₆ + 6NH₃ → 3[BH₂(NH₃)₂]⁺[BH₄]⁻ → (heat) 2B₃N₃H₆ + 12H₂.
• Structure: planar six-membered ring of alternate B and N atoms, both sp²; B–N bond length 144 pm (between single 154 and double 136 pm). N's lone pair is delocalised into B's empty p orbital, giving partial aromatic character.
• Isoelectronic and isostructural with benzene, but more reactive (polar B–N bonds; adds HCl readily).

Allotropes of carbon, carbides, silicates

Carbides

• Ionic (salt-like): formed by groups 1, 2, Al. Methanides (Be₂C, Al₄C₃ → give CH₄ with water), acetylides (CaC₂ → gives C₂H₂: CaC₂ + 2H₂O → Ca(OH)₂ + C₂H₂).
• Covalent: SiC (carborundum, abrasive), B₄C – very hard.
• Interstitial (metallic): C in holes of transition metal lattice, e.g. WC, TiC – hard, used in cutting tools.

Silicates – classification

Basic unit is the SiO₄⁴⁻ tetrahedron. Classified by how many O atoms are shared:

• Orthosilicates: discrete SiO₄⁴⁻ (zircon ZrSiO₄, forsterite).
• Pyrosilicates: Si₂O₇⁶⁻, one O shared (thortveitite).
• Cyclic silicates: (SiO₃)ₙ²ⁿ⁻ rings (beryl).
• Chain silicates: single chain (SiO₃)ₙ²ⁿ⁻ – pyroxenes; double chain (Si₄O₁₁)ₙ⁶ⁿ⁻ – amphiboles (asbestos).
• Sheet silicates: (Si₂O₅)ₙ²ⁿ⁻ – mica, talc, clay.
• Three-dimensional: SiO₂ (quartz), feldspars and zeolites (some Si replaced by Al).

Nitride of sulphur (S₄N₄) and fluorocarbons

• Preparation: 6S₂Cl₂ + 16NH₃ → S₄N₄ + S₈ + 12NH₄Cl.
• Structure of S₄N₄: an 8-membered cage (cradle shape). The 4 N atoms lie in a square plane; the 4 S atoms form a tetrahedron (two above and two below the plane). Two pairs of S atoms have weak transannular S····S interactions (≈ 258 pm). All S–N bonds are equal (≈ 162 pm) due to delocalisation. Orange solid, explodes on shock.

Fluorocarbons

Hydrocarbons in which H is replaced by F, e.g. CF₄, C₂F₆, Teflon (–CF₂–CF₂–)ₙ. Very stable due to the strong C–F bond; non-flammable and inert. Uses: Teflon non-stick coatings and insulation; Freons (CCl₂F₂) as refrigerants (now restricted, they damage ozone); perfluorocarbons as lubricants and blood substitutes.

Oxides and oxyacids of phosphorus

• P₄O₆ (written P₂O₃): based on the P₄ tetrahedron; one O atom is inserted into each of the 6 P–P edges. Each P is bonded to 3 O. P–O–P angle ≈ 127°.
• P₄O₁₀ (written P₂O₅): same as P₄O₆ plus one terminal O on each P (P=O, 143 pm, shorter than bridging P–O 160 pm, due to pπ–dπ bonding). Strong dehydrating agent.

In all, P is tetrahedral and has one P=O bond. Only H atoms on –OH are ionisable.

Peroxyacids of sulphur

• Caro's acid (peroxomonosulphuric acid, H₂SO₅): HO–SO₂–O–OH. Preparation: H₂SO₄ + H₂O₂ → H₂SO₅ + H₂O; or from chlorosulphonic acid: ClSO₃H + H₂O₂ → H₂SO₅ + HCl.
• Marshall's acid (peroxodisulphuric acid, H₂S₂O₈): HO–SO₂–O–O–SO₂–OH. Preparation: electrolysis of cold, concentrated H₂SO₄ at high current density: 2HSO₄⁻ → H₂S₂O₈ + 2e⁻ (at anode); or 2ClSO₃H + H₂O₂ → H₂S₂O₈ + 2HCl.
• In both, S is tetrahedral in +6 state; the –O–O– (peroxo) link makes them strong oxidising agents.

Interhalogens and polyhalides

Interhalogens are compounds of two different halogens: types AX, AX₃, AX₅, AX₇ (A = larger, less electronegative halogen).

• Preparation: direct combination. Cl₂ + F₂ → 2ClF (473 K); Cl₂ + 3F₂ → 2ClF₃ (573 K); I₂ + 5F₂ → 2IF₅; IF₅ + F₂ → IF₇.
• Properties: more reactive than halogens (A–X bond weaker than X–X, except F–F); hydrolyse to give halide and oxyhalide: ICl + H₂O → HCl + HOI. ClF₃ and BrF₃ are strong fluorinating agents.
• Structures (VSEPR): ClF₃, BrF₃ – T-shaped (sp³d); BrF₅, IF₅ – square pyramidal (sp³d²); IF₇ – pentagonal bipyramidal (sp³d³).

Polyhalides

Halide ions combine with halogens or interhalogens to give polyhalide ions: KI + I₂ → KI₃. Classified as homo-polyhalides (I₃⁻, I₅⁻) and hetero-polyhalides (ICl₂⁻, ICl₄⁻, IBr₂⁻).

• I₃⁻: central I has 2 bond pairs + 3 lone pairs (sp³d) → linear.
• I₅⁻: V-shaped (bent) chain, central I⁻ joined to two I₂ units [I–I····I⁻····I–I], angle ≈ 95°.
• ICl₄⁻: 4 bond pairs + 2 lone pairs (sp³d²) → square planar.

Likely exam questions

Long answer

1. Describe the structure and bonding in diborane. How is borazine prepared? Give its structure.
2. Give the structures of H₃PO₂, H₃PO₃, H₃PO₄, H₄P₂O₆ and H₄P₂O₇. Comment on their basicity and reducing nature.

Short answer

1. Compare the structures of graphite and diamond.
2. How are Caro's acid and Marshall's acid prepared? Give their structures.
3. What are polyhalides? Explain the structures of I₃⁻ and ICl₄⁻.